Sulphur Dioxide, SO2

Occurrence

The gases issuing from volcanoes contain sulphur dioxide in the free condition, and the gas is also present dissolved in the neighbouring springs. Sulphur dioxide is present in minute quantities in the air of towns, especially such as are occupied on an extensive scale with industries consuming much coal. The air of tunnels through which steam locomotives pass is frequently highly charged with sulphur dioxide.

On account of its characteristic odour, sulphur dioxide was early recognised as a product of the combustion of sulphur, although the gas was not isolated until Priestley in 1775, by means of his mercury trough, collected it from the interaction of mercury and sulphuric acid. Lavoisier almost immediately afterwards showed that the new substance was a compound of sulphur and oxygen inferior to sulphuric acid in oxygen content.

Sulphur Dioxide Preparation

From Sulphur

The simplest method of producing sulphur dioxide consists in burning sulphur in air or oxygen; the element ignites at about 250° to 260° C. The primary product, however, is not perfectly pure dioxide; the product of combustion in oxygen contains 2 or 3 per cent, of sulphur trioxide; in air the proportion of trioxide is still higher and may attain to 7 per cent, of the sulphur burnt, the nitrogen apparently acting as catalyst. Increase of pressure favours the production of the trioxide, as its formation involves a greater volume reduction, and experiment shows that under a pressure of 40 to 50 atmospheres of oxygen almost 50 per cent, of the sulphur may be converted into trioxide.

The necessary oxygen need not be obtained from the air. Sulphur dioxide is readily obtained by heating sulphur with certain oxides, dioxides or peroxides. Sulphur vapour and steam heated at a suitable temperature in the presence of iron oxide as catalyst react to form sulphur dioxide and hydrogen which may be recovered separately. With manganese dioxide and powdered sulphur the reaction proceeds as follows:

MnO2 + 2S = MnS + SO2,

manganese sulphide constituting the residue.

Similarly, when sulphur is heated with 2 to 3 times its weight of dehydrated ferrous sulphate or copper sulphate, sulphur dioxide is produced together with the sulphide of the metal:

FeSO4 + 2S = FeS + 2SO2.

From Sulphides

By heating metallic sulphides in air sulphur dioxide is obtained in a more or less pure condition. In the commercial preparation of sulphuric acid, for example, iron pyrites, FeS2, usually containing copper pyrites, Cu2S.Fe2S3, and contaminated with varying amounts of siliceous gangue, arsenic and other compounds, is roasted in air to yield the necessary sulphur dioxide. The main reaction taking place may be represented by the equation:

4FeS2 + 11O2 = 2Fe2O3 + 8SO2,

but in addition to the presence of volatile impurities from the pyrites such as arsenious oxide, sulphur trioxide is also present, for the metallic oxides produced during the roasting process exert a catalytic effect, thereby facilitating the further oxidation of the dioxide.

The use of zinc blende and mixed zinc ores as a convenient source of sulphur dioxide for sulphuric acid manufacture has been advocated. The residue from the roasting is often more valuable than in the case of pyrites, but the roasting process is more troublesome and the presence of fluorspar in the blende exerts a deteriorating influence on the lead chambers.

On roasting lead sulphide (galena) with lime, sulphur dioxide is evolved. It is assumed that the reactions which take place are as follows:

When hydrogen sulphide is burned in excess of air, sulphur dioxide is formed:

2H2S + 8O2 = 2H2O + 2SO2.

This reaction has been made the basis of a patent for the commercial production of the gas.

From Sulphuric Acid

Concentrated sulphuric acid can be made to yield sulphur dioxide by a variety of processes:

If merely raised to a red heat, as, for example, when it is dropped into a red-hot retort, sulphuric acid undergoes decomposition according to the equation:

2H2SO4 = 2H2O + 2SO2 + O2,

and the sulphur dioxide can be separated by passing the dried gaseous product through a vessel immersed in a freezing mixture. This method was formerly used in preparing a mixture of sulphur dioxide and oxygen for the manufacture of sulphur trioxide.

Sulphuric acid, when heated with copper or mercury, is reduced to sulphur dioxide. With the former metal, cupric sulphate is formed together with cuprous sulphide and even cuprous sulphate. The main reaction at a temperature of 130° to 170° C. appears to be:

6Cu + 6H2SO4 = 4CuSO4 + Cu2S + SO2 + 6H2O.

With mercury, the mercuric sulphate may be accompanied by mercurous sulphate.

Under suitable conditions sulphur dioxide may be obtained by the action of sulphuric acid upon many other metals. Thus, finely divided iron decomposes acid of density 1.75 at 200° C., yielding the gas.

Sulphuric acid is reduced to sulphur dioxide and water by prolonged contact with hydrogen at ordinary temperatures. The reaction is fairly rapid at 250° C.

Carbon and sulphur also reduce sulphuric acid, the former giving a mixture containing carbon and sulphur dioxides, whilst the latter yields a steady stream of pure sulphur dioxide. In the laboratory preparation of the gas from sulphuric acid and sulphur, however, it is necessary to have a wide delivery tube to prevent the exit being choked by the sulphur. The reactions are:

C + 2H2SO4 = CO2 + 2H2O + 2SO2. S + 2H2SO4 = 2H2O + 3SO2.

Both of the foregoing methods have been used commercially in preparing sulphur dioxide.

From Salts

Sulphites, thiosulphates and polythionates readily liberate sulphur dioxide when treated with an acid, or, in the case of the salts of the heavier metals, when merely heated. Calcium sulphite mixed with plaster of Paris and moulded into cubes forms a suitable material for use with sulphuric acid in an automatic gas generator such as Kipp's apparatus.

(b) When the sulphates of the metals of the alkaline earths are heated in the presence of a suitable metal such as iron, or the lower oxides of certain metals, or carbon, sulphur dioxide is obtained. The sulphates are reduced to the corresponding sulphides. The reduction is complete in half an hour at 750° C. for calcium sulphate, at 850° C. for strontium sulphate and at 950° C. for barium sulphate. At temperatures of about 150° C. higher and with an insufficiency of iron, a rapid evolution of sulphur dioxide takes place. The mixture 16CaSO4 + 15Fe gives an 80 per cent, theoretical yield of sulphur dioxide. As a result of these reactions certain well-defined ferrites are formed, for example, Ca3Fe4O9, Sr2Fe2O5 and BaFe2O4. Gypsum heated at 1400° C. with blast furnace slag containing calcium sulphide gives a high yield of sulphur dioxide:

CaS + 3CaSO4 = 4CaO + 4SO2.

Several other processes have been devised for utilising gypsum and anhydrite as sources of sulphur dioxide; for example, the mineral may be heated at 800° C. with coke, coal, or powdered shale and clay in a slightly oxidising atmosphere, when cement remains as a marketable residue, and the exit gases contain 6 to 7 per cent, of sulphur dioxide.

Physical Properties

Sulphur dioxide is a colourless gas possessing a pungent, choking odour, and exerting an extremely irritating action on the mucous membrane. In consequence of this it has been used in poison gas warfare, but it is not efficient against well-equipped troops on account of the ease with which it is extracted from the air by respirators. Approximately 0.0005 per cent, can be detected by the sense of smell, whilst 0.05 per cent, is unendurable. Vegetation is injured if the concentration of the gas exceeds 0.003 per cent.

The relation between the weight of the gas at a pressure p, namely Wp, and the weight at 760 mm. (W760) is given by the expression:

The critical temperature and pressure are, respectively, 157.5° C. and 77.8 atmospheres. At 0° C. the coefficient of expansion is 0.003978, and between 0° and 20° C. 0.00396, a value that is above the value for permanent gases. At higher temperatures, however, the coefficient becomes smaller, and between 400° and 1700° C. possesses the normal value. Between these two temperatures, therefore, there must be an absence of any appreciable decomposition or dissociation. The degree of dissociation into sulphur and oxygen at 1500° C. and 1 atmosphere pressure is 5.9×10-5, whilst under a pressure of 0.01 atmosphere it is 27×10-5.

The refractive index for sodium light is 1.00061 at 0° C. and 760 mm. The electrical conductivity is exceedingly small. The specific heat of sulphur dioxide is 0.1544 (water = 1), and the ratio = Cp/Cv = 1.290, a value agreeing quite well with that usual for tri-atomic gases, namely 1.30. Measurement has also been made of the gaseous viscosity, which is 1.253×10-4 C.G.S. units at 18.0° C. and 1.630×10-4 at 100° C. Investigation of the dielectric constant for gaseous sulphur dioxide is rendered difficult by the tendency of the gas to decompose under electrical stress. The value at 0° C. and atmospheric pressure has been found to be 1.00993. The electric moment of. the gas, calculated from Debye's equation, has the value (in C.G.S. E.S.U. ×1018) of 1.611.

The ultra-violet absorption spectra of the gas and its aqueous solutions have been investigated. In the case of the gas the most distinct absorption band lies between 317.9 and 263.0 μμ, whilst solutions of the gas exhibit a characteristic band at 276 μμ. Absorption of the infra-red by sulphur dioxide has also been studied.

Sulphur dioxide is diamagnetic; at 0° C. and 760 mm. pressure the magnetic susceptibility is -8.5×10-10.

The gas is readily soluble in water, the process of solution being accompanied by an evolution of heat. The heats of solution for various concentrations have been measured at 25° C., and the relation between -ΔH, the total heat of solution per gram-molecule of SO2 dissolved, and N, the number of molecules of water to one molecule of SO2, is found to be

-ΔH = 4911.6 + 1105.26 log N.

The following table gives typical experimental data:

Total heat of solutions of sulphur dioxide

Δt, ° C.

Grams SO2.

Grams H2O.

Mols. H2O/Mols. SO2'

-ΔH Obs., cals./mol.

-ΔH, Calculated from Equation, cals./mol.

0.180

0.2712

163.28

2140.9

8587

8592

0.260

0.3936

160.12

1446.6

8412

8405

0.447

0.7448

172.59

824.01

8122

8134

0.658

1.0973

168.18

545.01

7946

7936

0.947

1.6411

170.83

370.15

7745

7550

1.768

3.1480

167.59

189.30

7421

7428

4.190

7.9602

168.76

75.29

6994

6986

Below 50° C. the solubility of sulphur dioxide in water does not obey Henry's Law. In the following table the solubility is given as grams of sulphur dioxide per 100 grams of water when the partial pressure of the gas is equal to 760 mm. of mercury.

Solubility of sulphur dioxide in water

Temperature, ° C

Grams SO2 per 100 grams H2O.

(a)

(b)

10

15.4

15.39

15

12.55

12.73

20

10.4

10.64

25

8.95

8.98

30

7.8

7.56

40

5.8

5.54

50

4.5

4.14

60

. . .

3.24

At 15° C. and 760 mm., 1 c.c. of water absorbs approximately 44 c.c. of the gas.

The density of the saturated solution is as follows: at 0° C., 1.061; at 10° C., 1.055; at 20° C., 1.024. The following determinations for solutions of various concentrations have been made at 15.5° C.:

SO2 (per cent.)

0.99

2.05

2.87

4.04

4.99/td>

5.89

D15.54

1.0041

1.0092

1.0138

1.0194

1.0242

1.0287

SO2 (per cent.)

7.01

8.08

8.68

9.80

10.75

D15.54

1.0343

1.0389

1.0428

1.0482

1.0530

The aqueous solutions decompose slowly at ordinary temperatures, depositing sulphur, sulphuric acid being formed in solution. At 160° C. the decomposition is rapid.

Under suitable conditions solid products of variable composition may be obtained from aqueous solutions of sulphur dioxide, and hydrates containing 6 to 15 molecules of water have been described. It has been shown by a more recent investigation, however, that only one hydrate, namely SO2.6H2O, is formed. This gives a eutectic with ice at -2.6° C. The so-called higher hydrates consist of mixtures of the hexahydrate with ice.

Certain organic liquids also dissolve sulphur dioxide, and a direct comparison of the solubility in water and in chloroform has been effected by "partition" experiments, in which measurement is made of the distribution of the gas between the two solvents in contact with one another.

Alcohol is superior to water as a solvent for sulphur dioxide, absorbing more than 200 times its own volume of the gas at 0° C. and 760 mm.; thus, 1 c.c. of alcohol absorbs the following volumes:

Temperature, °C.

Volume of Gas dissolved, c.c.

0

216

5

175

10

142

15

116

20

96

25

84

Acetic acid under similar conditions dissolves more than 300 volumes or almost its own weight of the gas. Acetone dissolves about twice its weight or nearly 600 times its volume of sulphur dioxide, whilst camphor also absorbs more than 300 times its volume, forming a liquid solution. In the last two cases chemical combination undoubtedly occurs; the freezing-point curve of camphor-sulphur dioxide mixtures indicates the formation of two unstable compounds, namely, C10H10O.2SO2, m.pt. -45° C., and C10H10O.SO2, m.pt. -24° C.; these are probably active in the preparation of sulphuryl chloride in the presence of camphor. The use of methylcyclohexanone has been recommended as an absorbing liquid for the recovery of sulphur dioxide from waste gases.

Sulphur dioxide dissolves in aqueous solutions of inorganic salts frequently more readily than in pure water. With most salts, excluding sulphates, compounds appear to be formed in solution of the general type MX.SO2, where M and X stand for univalent metal and negative radical, respectively. The solubility curve of sulphur dioxide in sulphuric acid of concentration ranging from 55 to 98.5 per cent, is interesting. A minimum occurs at 85.8 per cent, acid, and from that point the curve inclines sharply upwards for both increase or decrease of sulphuric acid concentration.

Solubility of sulphur dioxide in sulphuric acid at 20° C and 760 mm.

Sulphuric Acid,

Grams SO2 in 100 grams of Acid.

55.1

5.13

68.9

4.16

80.2

3.12

82.5

2.99

84.2

2.88

85.3

2.83

85.8

2.80

86.5

2.82

88.1

2.90

90.8

3.10

92.8

3.21

94.6

3.50

96.5

3.83

98.5

4.03

At the point of minimum solubility of the sulphur dioxide the composition of the solvent closely approximates to that required for the monohydrate H2SO4.H2O (namely 84.5 per cent, acid), and it is significant that other physical properties of the acid pass through critical values at this concentration.

Sulphuryl chloride, SO2Cl2, is a still better solvent for sulphur dioxide and can absorb 187 volumes of the gas at the ordinary temperature.

Wood charcoal at 0° C. readily absorbs gaseous sulphur dioxide up to a volume even exceeding 100 times that of the charcoal; platinum black absorbs the gas to a relatively slight extent. The gas is also absorbed very appreciably by caoutchouc, the process merely being one of solution. The adsorption of sulphur dioxide by silicic acid gels has been measured. It is found that the maximum adsorption is exhibited by gels containing about 7 per cent, of water. In the absence of air the adsorption is reversible. Glass also adsorbs sulphur dioxide to a marked degree, depending on the time of contact, and it has been found impossible to remove the adsorbed gas completely from glass wool.

Solubility of Sulphur Dioxide in Sulphuric Acid at 20° C.

The formation of sulphur dioxide from rhombic sulphur and gaseous oxygen is exothermic to the extent of 71.08 Cals., the value being slightly higher, namely 71.72 Cals., for formation from monoclinic sulphur, the difference being due to the heat of transformation of a gram-atomic weight of sulphur from the rhombic to the monoclinic form.

Liquid Sulphur Dioxide

Sulphur dioxide was the first gas to be converted to the liquid state. It can be liquefied by passage through a tube cooled to below -10° C. in a freezing mixture, but commercially the liquid is produced by compression. The sulphurous gases from burning iron pyrites or some other suitable source, containing some 6 per cent, of sulphur dioxide, or certainly not less than 4 per cent., are extracted by cold water trickling down a tower and meeting the ascending gases; the solution thus obtained contains about 1 per cent, of the dioxide, and on warming, the dissolved dioxide is expelled. After being dried by passage through cooled pipes and finally through a chemical dehydrating agent such as sulphuric acid, the gas is liquefied by compression and the liquid enclosed in steel cylinders or, for smaller quantities, in glass cylinders. The liquid finds use in some ice-producing machines and also provides a convenient means of storing the substance.

Liquid sulphur dioxide is a colourless mobile fluid which boils at -10.02±0.1° C. at 760 mm.; the density varies considerably with the temperature, as indicated in the following table:

Temperature, °C.

Density.

Coefficient of Expansion.

-40

1.5331

0.00157

-20

1.4846

0.00164

0

1.4350

0.00175

2

1.3831

0.00192

40

1.3264

0.00223

60

1.2633

0.00261

80

1.1920

0.00315

100

1.1100

0.00390

The latent heat of evaporation per gram-molecule is 6150 calories, whilst the specific heat between -20° and + 10° C. is approximately 0.318. The vapour pressures at various temperatures from -20° to + 30° C. are given in the following table:

Temperature, °C.

- 20

- 10

0

+ 10

+ 20

+ 30

Vapour pressure (atm.)

0.65

1.04

1.58

2.34

3.35

4.67

The vapour pressure of sulphur dioxide at 50° C. is only about 8 atmospheres, although at 110° C. it exceeds 30 atmospheres, so that the sealed containing vessels are not likely to be submitted to very great strain so long as care is taken to prevent heating. The critical temperature is 157.50 ± 0.05° C. and the critical pressure 77.79 ± 0.05 atmospheres. The critical density is 0.5240 ± 0.0005. Measurement has also been made of the viscosity, the refractive index (n13°D = 1.350) and the compressibility.

The surface tensions of the liquid have been determined by the capillary tube method, and between -20° C. and 50° C. the values agree to within 5 per cent, with the equation:

γ (dynes/cm.) = 0.061534(157.5 – t)1.2,

and there appears to be no evidence of association of molecules over this range of temperature.

Although liquid sulphur dioxide is practically a non-conductor of electricity (its specific conductivity is 0.85×10-7 at 15° C.), it dissolves many substances giving solutions which conduct the electric current. Some salts, indeed, conduct better in liquid sulphur dioxide solution than in aqueous solution. The halides of the alkali metals, including ammonium and the alkylammoniums such as N(CH3)4Cl, produce solutions which, during electrolysis, yield the halogen at the anode, whilst sulphur is deposited on the cathode and a sulphite simultaneously formed. Cady and Taft, however, could detect no free sulphur in the cathode deposits obtained by electrolysis of solutions containing potassium iodide, iodate, ferricyanide or similar salt. In some cases the formation of additive compounds between solvent and solute has been detected, e.g. KI.4SO2, m.pt. +0.26° C.; KI.SO2, m.pt. -23.4° C. It has been suggested that the ionising power of sulphur dioxide may be connected with its high dielectric constant, 14.8 at 23° C.

Many organic substances are soluble in liquid sulphur dioxide, e.g. many alcohols, ether, resins, carbon disulphide, chloroform, benzene and alkaloids. It has been shown that under ordinary working conditions di-olefines are soluble and mono-olefines insoluble in liquid sulphur dioxide. This difference in solubility may be advantageously utilised in the refining of mineral lubricating oils, but it does not appear possible to separate naphthenes from paraffins by this method.

Phosphorus and sulphur are only sparingly soluble in the liquid, but sulphur monochloride is miscible. The binary systems formed by liquid sulphur dioxide with the tetrachlorides of carbon, tin and titanium, and with the tetrabromide of tin, have been investigated. The liquids are only partly miscible at lower temperatures and compound formation does not occur. With carbon tetrachloride, the critical solution temperature is -29.3° C., two liquid phases being possible down to -45° C.

The molecular ebullioscopic constant of sulphur dioxide is 15.0. When free from moisture, liquid sulphur dioxide is without action on iron and metals generally. The commercial liquid is generally quite free from sulphur trioxide.

Solid Sulphur Dioxide

Solid Sulphur Dioxide has been obtained by rapid evaporation of the liquid, part of which becomes cooled to the point of solidification. The solid is colourless and melts at -72.7° C.

Chemical Properties

Gaseous sulphur dioxide exhibits a tendency to undergo chemical change with - formation of an equilibrium mixture with sulphur trioxide and sulphur:

3SO2 ⇔ S + 2SO3.

Thus it becomes cloudy when exposed to strong illumination, and the presence of free sulphur in the gas at 1200° C. can be detected by Deville's "hot and cold tube" method. Light of any wave-length, within the absorption band, if it is of sufficient intensity, is capable of bringing about the decomposition of sulphur dioxide. The change does not take place, however, if the gas is absolutely dry. Slow decomposition as represented by the foregoing equation can also be effected by subjecting the gas to prolonged spark discharge. In all probability this conversion of sulphur dioxide into sulphur and sulphur trioxide is merely a special example of the power of gaseous sulphur dioxide at high temperatures to effect the oxidation of reducing agents such as hydrogen and carbon, the reducing agent or oxidisable substance in this case being part of the sulphur dioxide itself. At 2200° Abs. sulphur dioxide is not appreciably dissociated.

In the dry condition both gaseous and liquid sulphur dioxide are without action on dry litmus paper. Dry sulphur dioxide does not react with dry hydrogen at 100° C. or 280° C.

Sulphur Dioxide as an Oxidising Agent

Sulphur dioxide does not support the combustion of most substances which burn in oxygen, but many metals, e.g. sodium, potassium, magnesium and finely divided lead, when heated in a stream of the gas undergo conversion into a mixture of sulphide and oxide or sulphite, so much heat being liberated that the mass becomes incandescent; in the case of the alkali metals some thiosulphate also may be formed.

Molten copper absorbs several times its volume of sulphur dioxide, the product being a mixture of cuprous oxide and cuprous sulphide, which remains in solution in the metal. When dry, sulphur dioxide has no action on iron even at 100° C., but the metal is slightly attacked by the moist gas.

Hydrogen may be oxidised by sulphur dioxide under suitable conditions; thus when a mixture of the gases is passed over finely divided nickel or over nickel sulphide at a dull red heat, water, hydrogen sulphide and sulphur are produced. If the gases are passed sufficiently slowly the whole of the sulphur dioxide is decomposed. Cobalt sulphide and, with less efficiency, ferrous sulphide, may also be used as catalysts in this hydrogenation. A similar reaction occurs at a slightly higher temperature in the absence of a catalyst, but in this case the sulphur is obtained mainly in the free condition; the presence of a little water vapour facilitates the reaction.

Hydrogen iodide reduces dry sulphur dioxide to sulphur, but in the presence of moisture, the iodine thus produced tends to oxidise some of the remaining sulphur dioxide to sulphuric acid.

The interaction of hydrogen sulphide and sulphur dioxide to form water and sulphur, according to the equation

2H2S + SO2 ⇔ 2H2O + 3S,

does not take place in the absence of a liquid. Liquid hydrogen sulphide and liquid sulphur dioxide do not react even on boiling together. In the ordinary way the presence of a little water enables the interaction to take place, but many liquids other than water, for example, ethyl and methyl alcohols, glycerol and acetone, are able to bring about the reaction. There is no reaction between the gases at the ordinary temperature in the presence of phosphorus oxychloride, chlorobenzene, carbon tetrachloride, benzene, etc. There seems to be no connection between the dielectric constant and the catalytic activity. Sulphur dioxide and hydrogen sulphide do not react in the presence of the vapour of water or alcohol; it is essential that a liquid be present, probably because the liquid acts as a solvent for the gases. Alcohol is more effective than water, the gases being more soluble in alcohol.

At 100° C. there is no appreciable action between the moist gases, but sulphur is deposited on cooling. The reaction at higher temperatures has been investigated by passing the gases in equivalent proportions through a reaction vessel and analysing the rapidly cooled emergent gases. At 300° C. equilibrium is attained very slowly, the reaction taking place on the walls of the vessel. Using quartz powder as a surface catalyst the equilibrium is reached more rapidly, at 450° C., while above 600° C. such a catalyst is no longer necessary. The values for the equilibrium constant are given by:

ranging from 1.18 to 0.0062 between 450° and 600° C., and the heat of the reaction,

3S2(gas) + 4H2O(gas) = 4H2S + 2SO2,

is of the order of 32.0 kgm.-calories, whilst the value calculated from the reaction isochore is 28.0 kgm.-calories. The reaction between the mixed gases is also catalysed by active charcoal at lower temperatures, but above 600° C. some sulphur dioxide is reduced by the carbon, whilst above 800° C. carbon disulphide begins to form.

When sulphur dioxide alone is passed over carbon at a red heat the latter undergoes partial oxidation, the products being carbon monoxide, carbon oxysulphide and carbon disulphide; no oxysulphide is obtained at a white heat. Both carbon monoxide and methane are oxidised by sulphur dioxide at high temperatures with formation of sulphur. It has been suggested that volcanic sulphur may, in part, be formed by the reduction of sulphur dioxide by methane, carbon monoxide and hydrogen, all of which are emitted by volcanoes.

Other chemical processes in which sulphur dioxide functions as an oxidising agent include the interaction of sulphur dioxide with stannous chloride or titanous chloride in aqueous solution containing hydrochloric acid, when the salts are converted into the corresponding tetrachlorides; also the interaction of the gas with organo-magnesium compounds, when the organic sulphide is produced to some extent.

That sulphur dioxide can function either as an oxidising agent or as a reducing agent, according to the conditions, has been shown in the case of certain chlorides, sulphates and phosphates. The concentration of acid which is added is an important factor, for, by varying the amount present, the sulphur dioxide can be made either to oxidise or to reduce. The action of sulphur dioxide on the chlorides of mercury illustrates this diversity of action. Quantitative results are produced only under very specific conditions. A solution of mercuric chloride (1:80) saturated with sulphur dioxide at 70° to 80° C. and kept at that temperature for a considerable period reacts in accordance with the equation:

In the case of tin salts the action of sulphur dioxide is complex. In warm acid solution stannous sulphide is first precipitated and then yellow stannic sulphide mixed with sulphur. In highly concentrated acid solution, however, the sulphides are not precipitated,-but hydrogen sulphide is evolved:

SO2 + 6HCl + 3SnCl2 = 3SnCl2 + 2H2O + H2S.

In moderately acid solution the two tin salts and the sulphur dioxide compete for the hydrogen sulphide.

The reduction of ferric chloride solutions by sulphur dioxide is utilised in both qualitative and quantitative analysis, but reduction is not complete in the presence of considerable excess of hydrochloric or sulphuric acid. Experiments show that ferrous chloride is oxidised by sulphur dioxide in concentrated hydrochloric acid solution in accordance with the equation:

4FeCl2 + SO2 + 4HCl = 4FeCl3 + 2H2O + S.

Under the most favourable conditions the highest percentage of ferric iron obtained is 8.8. The degree of oxidation is independent of the initial concentration of total iron. It has also been shown that oxidation by sulphur dioxide at 95° C. does not occur in solutions of ferrous chloride unless there are present at least 165 grams of "free" HCl per litre.

Ferric chloride in concentrated hydrochloric acid solution is reduced by sulphur to a slight extent. Solutions containing more than 18.3 per cent, of ferric iron in the presence of 33 per cent, of hydrochloric acid are slowly reduced when a mixture containing equal quantities by weight of sulphur dioxide and hydrogen chloride is passed in at 115° C. From consideration of these facts it has been assumed that the reaction

4FeCl2 + SO2 + 4HCl ⇔ 4FeCl3 + 2H2O + S

is reversible.

The dependence of sulphur dioxide as an oxidising agent on a high concentration of hydrogen chloride has led to the suggestion that hydrogen chloride and sulphur dioxide first interact to a slight extent, forming thionyl chloride. Evidence in support of this is supplied by the following reactions with the mercaptans:

The fact that sulphur dioxide reduces most easily in a very dilute acid medium and oxidises most readily in a strongly acid medium may be correlated by explaining oxidation and reduction on an ionic basis, oxidation being represented by the surrender of positive charges and reduction by the transference of negative charges.

In aqueous solution sulphurous acid ionises into H•, HSO3' and SO3' ions. In this condition it acts as a reducing agent. Thus:

2Fe••• + SO3' + H2O = 2Fe•• + SO4 +

In strongly acid solution the concentration of SO3' ions is reduced and hence there is less tendency for reduction to take place.

As in the case of ferric chloride, cupric chloride is only incompletely reduced by sulphur dioxide in concentrated hydrochloric acid solution, but in aqueous solution this forms an excellent method for the preparation of cuprous chloride:

It has also been shown that in the presence of concentrated hydrochloric acid sulphur can reduce a hot solution of cupric chloride in accordance with the equation:

6CuCl2 + S + 4H2O = 6CuCl + 6HCl + H2SO4.

It is believed that the reaction

4CuCl + 4HCl + SO2 = 4CuCl2 + 2H2O + S

is reversible, but that the reversibility is normally obscured by the reaction between the sulphur and cupric chloride in the presence of the hydrochloric acid.

Sulphur dioxide and sulphurous gases attack basic rocks and glasses superficially at high temperatures (900° C.) with the formation of water-soluble sulphates, chiefly sodium sulphate. It is probable that such reaction and the solution of the products in hot springs during the early post-volcanic period explain the origin of alkaline sulphated thermal waters.

The capacity of sulphur dioxide to act either as an oxidising or as a reducing agent has been studied electrochemically in the case of solutions containing ferrous and ferric ions. Results show that an increase in acid concentration is accompanied by a rise in the sulphur dioxide potential, whilst under the same conditions there is a diminution in the ferric-ferrous iron potential. Experiments on the cathodic reduction of sulphur dioxide in acid solution show that hydrogen sulphide is not formed as a primary product.

Many sulphides, both natural and artificial, react with sulphur dioxide at high temperatures with formation of the corresponding sulphate and sulphur, thus:

RmSn + 2nSO2 = Rm(SO4)n + nS2.

In the case of the alkali and alkaline earth metals, the sulphate is the sole product; with copper, lead, bismuth and antimony, the product contains the metal, formed by reduction of the sulphate by unchanged sulphide. In cases where the sulphate is unstable at the temperature of reaction, as with zinc, cadmium, aluminium, tin, chromium, iron, cobalt and nickel, the oxide is the final product. The action may be catalytically accelerated by the addition of triferric tetroxide, Fe3O4.

Sulphur Dioxide as an Unsaturated Reagent

Sulphur dioxide exhibits a group of additive reactions indicative of its unsaturated nature.

Chlorine combines with the gas when the mixture is exposed to sunlight or when the sulphur dioxide is employed in solution in acetic acid or in camphor, the product being sulphuryl chloride; no combination occurs in the dark in the absence of a catalyst. In the presence of water the products are sulphuric acid and hydrochloric acid. Bromine does not combine in this manner, nor does hydrogen chloride, the freezing-point curve of the latter with sulphur dioxide showing only the formation of a eutectic mixture.

Sulphur dioxide combines with oxygen to form sulphur trioxide; reaction occurs when the gaseous mixture is passed over certain porous substances, such as sugar charcoal, finely divided platinum, or certain oxides, for example, iron, copper or cobalt oxide. The rate of reaction is practically independent of the concentration of oxygen, but is proportional to that of the sulphur dioxide and inversely proportional to that of the trioxide. The mode of action of such "contact substances" is not clearly understood, although it is probable that the effect is due, at least in part, to a primary condensation of one or both of the reagents on the catalyst. Traces of moisture are necessary to the reaction. Hargreave's process for the conversion of sodium chloride into sodium sulphate is an interesting application of "contact oxidation" of sulphur dioxide, in which the sodium chloride itself acts as the contact substance. A current of steam, sulphur dioxide and air is passed over the heated salt; if a little iron oxide or copper oxide be mixed with the salt, a quantitative conversion may be effected at 500° C.

Peroxides, dioxides and many basic oxides react, often vigorously, with sulphur dioxide, forming the corresponding sulphate as one of the products of the reaction, the actual change in each case varying with the nature of the metal. Sodium peroxide burns brilliantly when sprinkled into the gas. Lead dioxide becomes incandescent, whilst manganese dioxide reacts readily in the presence of water, the resulting solution containing dithionate and sulphate in addition to sulphite. Hydrogen peroxide gradually yields sulphuric acid.

Another indication of the unsaturated nature of sulphur dioxide is given by its action on the alkali iodides and thiocyanates. Complexes of the type MI.xSO2 are obtained, where M is an alkali metal.

Many aromatic hydrocarbons, for example, benzene, ethylbenzene, toluene, cymene and tetrahydronaphthalene, yield additive compounds. Such are also formed with liquid cyclic hydrocarbons in the absence of moisture and phenols, and use has been made of this fact to remove sulphur dioxide from a dry gas containing it. Additive compounds are also formed with methyl alcohol, thus CH3OH.SO2 and 2CH3OH.SO2, the existence of which has been demonstrated definitely by means of the freezing-point curve. The additive compound with camphor has already been mentioned.

Sulphur dioxide reacts with ammonia yielding amidosulphinic acid, NH2.SO2H, which, in the presence of excess of ammonia, is accompanied by the ammonium salt together with a red substance, triammonium imidodisulphinate, NH4.N(SO2.NH4)2.

Sulphur dioxide forms a substitution product in its reaction with phosphorus pentachloride, the products being thionyl chloride and phosphorus oxychloride:

PCl5 + SO2 = POCl3 + SOCl2.

Sulphur Dioxide as a Reducing Agent

When treated in a current of sulphur dioxide, nitrates are reduced, with formation of nitrous gases and sulphates; chlorates likewise are reduced, chlorine peroxide being obtained below 60° C., whilst above this temperature the volatile products are sulphur trioxide and chlorine.

Chromates and permanganates are vigorously reduced. Thus, on passing a current of sulphur dioxide through an acidified aqueous solution of potassium dichromate, reduction to chromium sulphate occurs, the orange solution becoming purple-green and depositing, upon concentration, crystals of potassium-chromium alum. Permanganates react similarly, the manganese being reduced to manganous sulphate and the solution becoming practically colourless. Thus:

2KMnO4 + 5SO2 + 2H2O = K2SO4 + 2MnSO4 + 2H2SO4.

This reaction is the basis of a useful volumetric method of estimating sulphites in solution.

Physiological Action

Sulphur dioxide exerts a decidedly toxic effect on plants and animals, and has been used in poison gas warfare; even as little as 0.04 per cent, by volume in the atmosphere will cause symptoms of poisoning in human beings after a few hours; in larger quantities, either gaseous or dissolved, the effect may be fatal. The gas acts as a direct blood poison and also affects the blood circulation. The sulphites are not poisonous.

Applications of Sulphur dioxide

Gaseous sulphur dioxide is used for bleaching certain natural colours, e.g. in wool, silk, feathers, straw and sugar; its action is less powerful than that of chlorine and some yellow flowers are unaffected by it. In some cases the bleaching, which occurs in the presence of moisture, is probably due to a reducing action, but for the most part the colour, for example, with red rose petals, can be restored partly by warming or treating with dilute sulphuric acid; the bleaching in the latter case is probably due, at least partially, to the formation of a feeble compound between the sulphur dioxide and the pigment.

The disinfectant action of sulphur dioxide was recognised by the ancient Greeks and is mentioned by Homer. To-day the gas is less favoured than in the past as a disinfectant although it is still largely used; it is also applied, more generally in the form of sulphite, for preservative purposes. Fresh whole fruit may be preserved in sealed vessels in an aqueous solution containing 0.08 to 0.1 per cent, of sulphur dioxide; absorption by the fruit occurs and the concentration falls to 0.04 per cent., but the growth of micro-organisms is inhibited, and jam made from such fruit is said to be of superior quality to that made from pulp. The application of sulphur to vines for the prevention of disease possibly owes its efficiency to a very slow formation of sulphur dioxide.

Liquid sulphur dioxide finds occasional use as a refrigerating liquid for the manufacture of ice and as a solvent for the extraction of fats and oils from bones and other waste animal matter. It is also employed in the refining of natural petroleum, owing to its property of dissolving aromatic and other heavy hydrocarbons which are present in petroleum distillates. When the distillate is shaken with liquid sulphur dioxide at a low temperature, separation into two layers occurs, one of which is the sulphur dioxide solution, the other the purified distillate containing paraffin hydrocarbons and naphthenes, which remain unaffected.

The gas is used in various chemical industries. For example, it forms an intermediate stage in the production of sulphuric acid and of sulphites; it is also applied in the preparation of chemical substances such as selenium, tellurium, quinol, etc. To some extent sulphur dioxide finds application for fire-prevention and extinguishing, and much is used in the form of alkali sulphite in the production of wood pulp for paper and artificial silk manufacture.

Detection and Estimation

The detection of sulphur dioxide, except by its odour and a few special reactions such as its reduction to hydrogen sulphide by hydrogen and its reducing action on potassium chromate paper, is generally effected by the same tests as are applied to the dissolved gas. The gas may also be recognised by lowering into it a rod which has been thinly coated with a layer of moist zinc nitroprusside, rendered transparent by exposure to ammonia. According to the amount of sulphur dioxide present the coating turns rose or deep red in colour. The addition of pure zinc to a solution of sulphurous acid causes reduction to hydrosulphite, which may be detected by its action on Methylene Blue.

Sulphur dioxide is conveniently estimated volumetrically by absorption in sodium hydroxide solution and, after acidifying with hydrochloric acid, titrating the solution with potassium iodate. Direct titration of sulphur dioxide solutions does not give consistent results except under special conditions. The solution should be made strongly alkaline and a small amount of sugar added; it is then put into a burette and run into a definite quantity of standard iodine solution made strongly acid with hydrochloric acid. For the accurate determination of small amounts of sulphur dioxide in gaseous mixtures the excess of iodine method is recommended. In this method the sulphur dioxide is dissolved directly in excess of iodine and the excess determined by means of thiosulphate and starch. Oxidation with permanganate under definite conditions is a usual method for estimation of the gas, permanganate solution being more stable than iodine solution. A method of estimating sulphur dioxide in a solution of bisulphite by the addition of mercuric chloride and subsequent titration with sodium hydroxide solution is described by Debucquet. For the estimation of sulphur dioxide in wines, the wine is heated for some time with excess of sodium hydroxide solution, and the excess alkali then neutralised with dilute sulphuric acid and the solution titrated with iodine and starch.